Boiling point

The boiling point of a substance is the temperature at which the vapour pressure of its liquid phase becomes equal to the external pressure. At this temperature, the liquid undergoes a phase transition and forms vapour bubbles within its bulk. Boiling points vary significantly with changes in external pressure, which is why liquids behave differently at various elevations or under controlled laboratory conditions. While water boils at 100 °C at sea level under one atmosphere of pressure, reduced pressure at higher altitudes results in a lower boiling point. Different liquids exhibit distinct boiling points under the same pressure because their vapour pressures differ at any given temperature.

Normal and Standard Boiling Points

The normal boiling point of a liquid is defined as the temperature at which its vapour pressure equals one atmosphere. Under these conditions, vapour bubbles can form within the liquid and rise to the surface. The standard boiling point, as defined by the International Union of Pure and Applied Chemistry (IUPAC) since 1982, is the temperature at which boiling occurs under a pressure of one bar (100 kPa). Although these values are close, they differ slightly because one atmosphere is equal to 101.325 kPa.
The boiling point is closely linked to the kinetic behaviour of molecules. When sufficient thermal energy is added, the vapour pressure of the liquid increases until it matches the surrounding pressure. At high elevations, where atmospheric pressure is reduced, a liquid reaches this condition at a lower temperature.

Evaporation and the Heat of Vaporisation

Liquids can transition into the vapour phase below their boiling points through evaporation. This is a surface phenomenon in which molecules at the liquid’s boundary escape into the surrounding air if they have enough kinetic energy to overcome intermolecular forces. Boiling, in contrast, occurs throughout the liquid as vapour bubbles form internally.
The heat of vaporisation denotes the thermal energy required to convert a specified amount of liquid into vapour at a constant pressure. This value varies among substances and reflects the strength of intermolecular forces. Substances with strong intermolecular attractions require more energy to vaporise and therefore exhibit higher boiling points.

Saturation Temperature and Pressure

A saturated liquid contains the maximum thermal energy possible without undergoing vapour formation, whereas a saturated vapour contains the minimum thermal energy it can retain without condensing. The saturation temperature corresponds to the boiling point at a specified saturation pressure. Adding heat to a saturated liquid at constant pressure causes boiling, whereas removing heat from a saturated vapour causes condensation.
Saturation pressure and saturation temperature vary directly. Increasing pressure raises the saturation temperature, while reducing pressure lowers it. At pressures above the critical point, liquid and vapour phases become indistinguishable, and boiling as a phase transition ceases. At pressures below the triple point, a distinct boiling point cannot exist.

Pressure Dependence and the Clausius–Clapeyron Equation

The boiling point is fundamentally dependent on external pressure. Increasing the surrounding pressure elevates the boiling temperature, while reducing the pressure lowers it. This relationship can be modelled using the Clausius–Clapeyron equation:
TB = 1 ⁄ [1 / T₀ – (R ln(P ⁄ P₀) ⁄ ΔHvap)] where TB is the boiling temperature at the pressure of interest, R the gas constant, P the vapour pressure at TB, P₀ a reference pressure, T₀ the boiling temperature at P₀, and ΔHvap the heat of vaporisation. This relation enables boiling points to be estimated when vapour pressure data are available.

Variations in Boiling Point: Elevation, Critical Point, and Triple Point

At higher elevations, atmospheric pressure is significantly lower, resulting in reduced boiling temperatures. For instance, water boils well below 100 °C at the summit of Mount Everest. Conversely, increasing the pressure raises the boiling point up to the critical point, beyond which the liquid and vapour phases merge into a supercritical fluid. The boiling point cannot be increased past this point.
Similarly, lowering the pressure decreases the boiling point until the triple point is reached. Below the triple point, a liquid phase cannot exist in equilibrium, and boiling is no longer defined.

Relationship Between Vapour Pressure and Normal Boiling Point

The normal boiling point of a substance is directly related to its vapour pressure curve. Liquids with high vapour pressures at a given temperature have lower normal boiling points, whereas those with low vapour pressures have higher boiling points. Vapour pressure charts displaying logarithmic relationships between temperature and vapour pressure clearly illustrate these trends.
For example:

• Methyl chloride, which has a relatively high vapour pressure at low temperatures, boils at –24.2 °C under atmospheric pressure.
• Liquids with lower vapour pressures require higher temperatures to reach one atmosphere and therefore have higher normal boiling points.

Boiling Point as a Physical Property

For pure substances, the normal boiling point is a characteristic physical property used for identification, classification, and comparison of volatility. A substance with a high normal boiling point is generally less volatile than one with a low boiling point. Many reference materials list both boiling and melting points as essential identifiers.
Certain compounds decompose before reaching their boiling points, especially large molecules such as polymers. In such cases, the boiling point may not be experimentally accessible. Between the triple point and critical point, stable compounds possess distinct boiling points dependent on pressure, while beyond the critical point only a single supercritical phase exists.
A compound’s behaviour at a given temperature can be inferred from its normal boiling point:

• Substances with boiling points below ambient temperature exist as gases at atmospheric pressure.
• Substances with boiling points above ambient temperature may exist as liquids or solids, depending on their melting points.

Factors Affecting Boiling Points

Boiling points are influenced by several structural and intermolecular factors:

• Ionic compounds typically exhibit very high boiling points due to strong electrostatic attractions.
• Metals often have high boiling points, although this varies widely with bonding structure.
• Covalent substances show increasing boiling points with greater molecular mass because of enhanced dispersion forces.
• Organic compounds display boiling point trends based on functional groups. Alcohols, carboxylic acids, and other strongly polar molecules often have higher boiling points due to hydrogen bonding, whereas alkenes, ethers, and halogenoalkanes show trends linked to mass and polarity.